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The Mole

The Mole. Chemistry 6.0. Getting to know the terms…. The following elements exist in nature as molecules: H 2 O 2 F 2 Br 2 I 2 N 2 Cl 2 S 8 P 4. MOLE RELATIONSHIPS. 1 Mole = 6.02x10 23 particles of substance (atoms, formula units, molecules)

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The Mole

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  1. The Mole Chemistry 6.0

  2. Getting to know the terms… The following elements exist in nature as molecules: H2O2 F2 Br2 I2 N2 Cl2S8 P4

  3. MOLE RELATIONSHIPS 1 Mole = 6.02x1023particles of substance (atoms, formula units, molecules) 1 Mole = mass (g) of substance from PT Also remember your formula information: 1 molecule = _________ atoms 1 formula unit = _________ ions or _________ atoms

  4. Mole ConversionsMUST use factor label! A. Moles & Mass 1. How many grams in 3.0 moles of water? know: 1 mole H2O = 2. How many moles in 60.0 g of copper? know: 1 mole Cu = B. Moles & Particles 1. How many atoms in 3.0 moles of copper? know: 1 mole Cu = 2. How many atoms in 3.00 moles of water? know: 1 mole H2O = know: 1 molecule H2O = 18.0 g H2O 54 g H2O 63.5 g Cu 0.945 g Cu 6.02 x 1023 atoms of copper 1.8 x 1024 atoms Cu 6.02 x 1023 molecules of H2O 3 atoms 5.42 x 1024 atoms

  5. Mole ConversionsMUST use factor label! C. Mass & Particles 1. How many atoms in 100.0 g of copper? know: 1 mole = _________ g copper 1 mole = 6.02 x 1023 __________ of copper 2. How many oxygen atoms are in 75.0 g of sucrose, C12H22O11? know: 1 mole = __________ g of C12H22O11 1 mole = 6.02 x 1023 _____________ of C12H22O11 1 molecule of C12H22O11 = 11 ________ of oxygen 63.5 atoms 9.480 x 1023 atoms Cu 342.0 molecules atoms 1.45 x 1024 atoms

  6. Molar Volume of Gases at STP Avogadro’s Law Amount - Volume Relationship. Equal volumesof gases at the same temperature and pressure contain an equal number of particles. volume constant 4 He 222 Rn molar mass 1 mole gas = 22.4 L = 6.02 x 1023 particles at STP (273 K & 1 atm)

  7. Therefore because of Avogadro’s Law if these three gases have the same number of particles and are at the same temperature and pressure, they must take up the same volume. He O2 Rn

  8. Molar Mass does not affect volume of a gas

  9. Avogadro’s Law • At STP, the amount of gas is directly proportional to the volume. Problem #1: Which of the following samples of gases occupies the largest volume, assuming that each sample is the same temp and pressure? 50.0 g Ne 50.0 g Ar 50.0 g Xe

  10. Avogadro’s Law V1 = V2 n1 n2

  11. n ____ P ____ (more gas, ____________) more collisions

  12. Ideal Gas Law Although no “ideal gas” exists, this law can be used to explain the behavior of real gases under ordinary conditions. P = pressure (atm) V = volume (L or dm3) n = number of moles R = 0.08206 L•atm/mol•K universal gas constant T = Kelvin temperature • Individual gas laws describe the relationships between these variables. • Ideal gas law relates all 4 variables that describe a gas at one set of conditions. PV = nRT

  13. Ideal Gas Law Problems • Calculate the volume of a gas balloon filled with 1.00 mole of helium when the pressure is 760. torr and the temperature is 0.oC. 22.4 L • Calculate the pressure, in atm, exerted by 54.0 g of xenon in a 1.00-L flask at 20.oC. 9.9 atm • Calculate the density of nitrogen dioxide, in g/L, at 1.24 atm and 50.oC. 2.16 g/L

  14. Empirical Formulas • Definition: always the smallest whole-number ratio of the atoms, or ions, in a formula • Use experimental data to find the empirical formula • Examples • Determine the empirical formula of a compound if a 2.500-g sample contains 0.900 g of calcium and 1.600 g of chlorine. • Determine the empirical formula for an iron oxide that is 78% iron. Name the compound. CaCl2 FeO iron(II) oxide

  15. Molecular Formula • Definition: the formula of a molecular compound. The molecular formula shows the actual number of atoms of each element present in 1 molecule of a compound. Molecular formula for benzene: C6H6 Empirical formula for benzene: • Molecular formula is always a whole-number multiple of the empirical formula. CH molecular formula = (empirical formula)n n = molar mass molecular formula molar mass empirical formula

  16. Example Find the molecular formula of a compound that contains 42.5 g of palladium and 0.80 g of hydrogen. The molar mass of the compound is 216.8 g/mol. Empirical formula - PdH2 Molecular formula – Pd2H4

  17. Concentration • Percent concentration by mass • (solute/solution) x 100% = % Concentration • Molarity (M) • Moles of solute/Liters of solution = mol/L • Molality (m) • Moles of solute/mass of solvent = mol/kg

  18. Molarity or Concentration a. Definition: number of moles of solute per liter of solution 1 L = 1 dm3 = 103mL = 103cm3 = 103cc b. Abbreviation: MUnits:mol/L c. Preparation of solutions • Need to know the desired volume & calculate the mass of needed solute. • Prepare 500. mL of 1.0 M NaCl Transfer ________ grams of NaCl to a 500-mL volumetric flask, and add water to the line. *Note: Always add acid to water. 29

  19. Problems • Calculate the molarity if 37 g of NaCl are dissolved in 150 mL of solution. • How many moles of HCl are present in 145 mL of a 2.25 M HCl solution? • How many grams of NaCl are contained in 2.5 L of a 1.5 M solution? 4.2 M NaCl 0.326 mol HCl 220 g NaCl

  20. Example Problems • Calculate the molarity of a solution that contains 8.50 g of calcium nitrate in 2.0 L. • Calculate the molality of a solution that contains 8.50 g of calcium nitrate in 125 g H2O.

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